1. Field of the Invention
This invention relates to the field of electrochemical batteries of both the primary and secondary (storage or rechargeable) type and, more particularly, relates to the field of electrochemical batteries having non-aqueous electrolyte systems.
2. Description of the Prior Art
In recent years, numerous proposals have been made for increasing the gravimetric and volumetric energy densities of electrochemical batteries through the application of highly reactive metals, e.g., the alkali metals and the alkaline earth metals, as anodic materials. Lithium metal has received by far the most attention in this regard due to its very low atomic weight and its being the most electronegative of all the metals. Batteries containing lithium or other light metal anodes cannot employ aqueous and other active hydrogen-containing electrolytes since contact of these metals with such electrolytes would result in oxidation of the latter and evolution of hydrogen gas. Accordingly, batteries of this type which are intended for service at ambient temperatures are provided with non-aqueous electrolytes in which electrically conductive salts are dissolved in organic aprotic solvents. Among the numerous electrically conductive salts which have heretofore been employed in non-aqueous electrolyte systems are the light metal and ammonium salts of such anions as the halides, halates, perhalates, haloaluminates, haloarsenates, halophosphates, haloacetates, phosphates, thiocyanates, sulfides, sulfates, cyanides, picrates, acetylacetonates, fluoborates, hydrides, borohydrides, and so forth. These and other electrically conductive salts have been dissolved in a wide variety of organic aprotic solvents including the normal and branched paraffins and cycloparaffins; aromatic hydrocarbons such as benzene, toluene and xylene; Lewis bases such as the tertiary amines; amides and substituted amides such as formamide; nitriles such as acetonitrile, propionitrile and benzonitrile; open chain and cyclic esters such as propylene carbonate, alkyl acylates and butyrolactone; oxysulfur compounds such as dimethylsulfoxide, dimethylsulfite and tetramethylene sulfone; and, open chain and cyclic ethers such as the poly (alkyleneoxy) glycols, dioxane and the substituted dioxanes, dioxolane, tetrahydrofuran and tetrahydropyran. Illustrative of primary and/or secondary batteries having light metal anodes and non-aqueous electrolytes are those described in U.S. Pat. Nos. 3,185,590; 3,393,092; 3,404,042; 3,413,154; 3,489,611; 3,531,328; 3,533,853; 3,542,601; 3,542,602; 3,578,500; 3,764,385; 3,918,988; 3,920,477; 3,928,067; 3,928,070; 3,953,302; and 3,982,958.
Ideally, the organic aprotic solvent selected for use in non-aqueous electrolytes should combine good solvency for the electrically conductive solute (to permit high levels of ionic conductivity) with long term stability in contact with the anode. In practice, the properties are inherently opposed to each other. The electrolyte solvents of high solvency, i.e., the polar solvents, are the least stable in contact with the highly electronegative light metals. The solvents of high stability, i.e., non-polar solvents such as the aromatic hydrocarbons, have the poorest solvency for the electrically conductive salts. One possible mechanism which would explain the poor stability of a polar aprotic solvent toward light metals is that the cation of the electrically conductive salt dissolved in the solvent behaves as a Lewis acid catalyst through association with the electronegative element of the solvent. Such association is thought to result in a shift of negative charge to the cation rendering the electronegative element and the carbn atom adjacent to it more positive. This in turn facilitates the electron transfer to the solvent molecule and the production of an anion radical as the initial step in the solvent decomposition process. Regardless of the precise nature by which the polar aprotic solvents are eventually degraded, the fact remains that the usefulness of these solvents in high energy battery systems is limited by their tendency to undergo decomposition in the presence of light metal anodes.